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The Atom and Its Properties Chapter 4 – Nucleus Chapter 5 – Electron Configuration Chapter 4 Objectives • Describe an atom’s structure and differentiate among the particles that make it up. • Identify the numbers associated with elements and explain their meaning . • Realize that the number of protons in a nucleus defines an element. • Calculate the average atomic mass given isotopes and relative abundance 2 Chapter 4 Vocabulary Chapter 4.3 Chapter 4.1 • Atomic Number • Dalton’s Atomic • Isotope Theory • Mass Number • Atom • AMU (Atomic Mass Unit) Chapter 4.2 • Average Atomic Mass • Electron • Nucleus • Proton • Neutron 3 Modern View of the Atom The nucleus is where the protons and neutrons are located and contain most of the atom’s mass. 4 Protons, Neutrons and Electrons 5 Particle Symbol Charge Electron Proton Neutron ep+ n0 11+ 0 Relative Mass 1/1840 1 ~1 •Sometimes Atomic Symbols are Displayed as: 6 Isotope Examples Isotope Atomic Mass Number Number Number Number Number Protons Neutron Electron Cr-52 24 52 33 222 7 42 86 What’s all this amu business? • • • 8 To simplify a system of indicating atomic masses since protons and neutrons have such extremely small masses, scientists have assigned the carbon-12 atom a mass of exactly 12 atomic mass units. (amu) The mass of 1 amu (1/12 the mass of carbon12) is very nearly equal to the mass of a single proton or neutron but not the same. 1 amu = 1.66 x 10-24 grams •Isotopes and Mass Number • Your text (p. 119) shows how to calculate the mass number for Cl given the % abundance of the isotopes. • Let’s do this for another element: Li • 6Li • 7Li is 7.59 % abundant; 6.015 amu is 92.41% abundant; 7.015 amu • Method 1: Use percentages. Think of this as a sample of 100 atoms. 9 •In Tabular Form Species Mass (amu) 6Li 6.015 (isotope) 7Li 7.015 (isotope) Li (100 atoms) Li (atom) 10 Abundance Mass x Abundance % (Weighted share) 7.59 45.65 92.41 648.26 100.00 694.41 6.94 Average Atomic Mass • The average mass of an atom is found by weighting the natural abundances of its isotopes. • Lithium • 6Li • 7Li (Method 2): Change % to fraction. 6.015 amu 7.59% = 0.0759 7.015 amu 92.41% = 0.9241 Mass (amu) Frac abund Mass share Avg mass = 6.015 amu x 0.0759 = 0.46 amu 7.015 amu x 0.9241 = 6.48 amu 6.94 amu/atom 11 •Displayed on Periodic Table 12 13 Electrons in Atoms Chapter 5 •Chapter 5 Objectives Compare wave and particle matters of light • See how frequency of light emitted by an atom is unique to that atom • Compare and contrast the Bohr and quantum mechanical models of the atom • Express the arrangements of electrons in atoms through orbital notations, electron configurations, and electron dot structures • 15 Chapter 5 Vocabulary Chapter 5.1 • Electromagnetic Radiation • Wavelength • Frequency • Amplitude • Electromagnetic Spectrum • Quantum • Photoelectric Effect • Photon • Atomic Emission Spectrum 16 Chapter 5.2 • Ground State • Quantum Number • Quantum Mechanical Model of the Atom • Atomic Orbital • Principal Quantum Number • Principal Energy Level • Energy Sublevel •Wave Nature of Light • Electromagnetic radiation displays wavelike behavior as it travels through space • Waves can be described by several common characteristics 17 •Characteristics of a Wave • • Waves transfer energy Properties of waves: • Frequency (ν – pronounced ‘nu’) Number of vibrations per unit time – Hz (cycle/second) • Wavelength (λ) Distance between points on two consecutive waves • Speed of wave is Frequency x wavelength Speed = ν x λ Frequency is the number of waves that hit this point in one second. Amplitude λ 1.5 1 0.5 0 0 200 400 600 -0.5 -1 -1.5 18 800 1000 1200 •Electromagnetic Spectrum Note: All EM Radiation travels at 3.00 x 108 m/s 19 •Electromagnetic Spectrum The speed of light (3.00 108 m/s) is the product of it’s wavelength and frequency c = λν. 20 The Electromagnetic Spectrum – all light is energy 21 •Electromagnetic Spectrum • Gamma Rays – Highest frequency, shortest wavelength. Can pass through most substances • X Rays – Lower frequency than Gamma rays. Can pass through soft body tissue but can’t pass through bone. • Ultraviolet (UV) Rays – Part of sunlight that causes sunburn 22 •Electromagnetic Spectrum • Visible Light – Sensitive to our eyes. Allows us to see color • Infrared – Less energy and longer wavelength than visible light. Felt as heat given off a heater or near a fire • Radio Waves – Lowest frequencies on the EM spectrum. Used by radio and over-the-air TV. 23 An electromagnetic wave has a frequency of 6.0 x 104 Hz. Convert this frequency into its corresponding wavelength. Which region of the EM spectrum does this correspond to? lxn=c l = c/n l = 3.00 x 108 m/s 6.0 x 104 /s l = 5.0 x 103 m It’s a radio wave (~103 meters) 24 7.1 •Practice Problems What is the frequency of green light, which has a wavelength of 520 nm. • A radio station broadcasts at 94.7 MHz. What is the wavelength of the broadcast? • 25 Answers •Nature of Light • Max Planck (18581947) studied the different light emitted from heated objects • Matter can only gain or lose energy in small specific amounts 26 •Nature of Light • A quantum is the minimum amount of energy that can be gained or lost by an atom • The energy of EM radiation is proportional to its frequency (E α ν) 27 •Photons Albert Einstein (18791955) proposed that while a beam of light had wavelike characteristics, it also can be thought of as a stream of tiny particles (or bundles of energy) called photons • Each photon carries a quantum of energy • 28 •Particle Nature of Light The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it. 29 When copper is bombarded with high-energy electrons, X rays are emitted. Calculate the energy (in Joules) of the X-rays if their frequency is 1.95 x 1018 Hz. E=hxn E = 6.63 x 10-34 (J•s) x 1.95 x 1018 /s E = 1.29 x 10 -15 J 30 7.2 Ch. 5.2 – Quantum Theory of the Atom Each element has only certain specific frequencies of light that are emitted when atoms absorb energy and become excited Where do we see this? fireworks neon signs stars 31 •Hydrogen Spectrum 32 •Balmer Plot • In 1885, Johann Balmer observed the lines of the spectrum fit this surprisingly simple formula: 1 1 RH 2 2 l n1 n2 1 • Where n1 =2 and n2 = 3, 4, 5, etc. 33 1/Labmda, m-1 •Balmer Plot 2500000 2400000 y = 1.0972E+07x + 4.0238E+02 2300000 R² = 1.0000E+00 2200000 2100000 2000000 1900000 1800000 1700000 1600000 1500000 0.12 0.14 0.16 0.18 0.2 0.22 0.24 1/2^2 - 1/n^2 RH is the slope of this line, 1.0972 x 107 m-1 34 •Electronic Energy Transitions • Neils Bohr (18851962) proposed the model the hydrogen atom (1913) to explain the discreet nature of the hydrogen spectrum. 35 •Electronic Energy Transitions • Neils Bohr’s model the atom (1913) Electrons exist only in discrete, “allowable” energy levels • Energy is involved in moving electrons from one energy level to another • Principal quantum number (n) - specifies the electron’s major energy level • The lowest energy is n=1, the next lowest in n=2, etc. • 36 •Bohr’s Model of the Atom (cont’d) Bohr suggested that an electron moves around the nucleus in only certain allowed circular orbits. n=2 n=1 37 •Energy Absorption/Emission 38 39 •Atomic Emission Spectra 40 •Origin of Line Spectra Balmer series 41 42 •Quantum or Wave Mechanics Schrodinger applied idea of ebehaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION E. Schrodinger 1887-1961 43 Solution gives set of math expressions called WAVE FUNCTIONS. Treated electrons as wavelike particles that became the Quantum Mechanical Model of the Atom. Waves •Wave motion: wave length and nodes “Quantization” in a standing wave 44 •Hydrogen Atom Solution Where: a0 is the Bohr Radius given by a0 = 4πεoh2/me2 Generalized Laguerre Polynomial m here is quantum number Constant = 2.18 x 10-18 J m is mass of electron 45 46 •Atomic Orbitals-Hydrogen 47 •Orbitals No more than 2 e- assigned to an orbital • Orbitals grouped in s, p, d (and f) sublevels • s orbitals p orbitals d orbitals 48 s orbitals p orbitals d orbitals s orbitals No. orbs. p orbitals d orbitals 1 3 5 2 6 10 No. e- 49 Energy Levels and Sublevels • Sublevels are grouped in energy level. • Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n which indicates relative size and energy of the orbitals • Row on PT indicates n 50 •Energy Levels and Sublevels n=1 n=2 n=3 n=4 51 •QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of quantum numbers: n (principal) l (angular) • Note: Energy Level sublevel (s, p, d, or f) which is its shape There are other quantum numbers that we will NOT discuss in detail. The ‘n’ and ‘l’ are sufficient. 52 QUANTUM NUMBERS Symbol Values n (principal) 1, 2, 3, .. l (angular) More commonly noted as: 53 Description Orbital size and energy level 0, 1, 2, 3, … n-1 Orbital shape s, p, d, f, …n-1 or type (energy sublevel) Levels and Sublevels When n = 1, then l = 0 Therefore, in n = 1, there is just 1 type of sublevel and that sublevel has a single orbital This sublevel is labeled s (“ess”) Each level has 1 orbital labeled s, and it is SPHERICAL in shape. 54 •Types of Atomic Orbitals 55 Types of Atomic Orbitals 56 •s Orbitals— Always Spherical Dot picture of electron cloud in 1s orbital. Surface density 4πr2y versus distance See Active Figure 6.13 57 Surface of 90% probability sphere •p Orbitals The three p orbitals lie 90o apart in space 58 •2px Orbital 59 3px Orbital •Hydrogen-like Orbitals (at most two electrons/orbital) n 1 2 3 4 60 Sublevel (l) s s p s p d s p d f Orbitals Max. Orbital Max Elec n2 2n2 s 1 2 s px, py, pz 4 8 s px, py, pz dxy,dxz,dyz,dx2-y2, dz2 9 18 s px, py, pz dxy,dxz,dyz,dx2-y2, dz2 And 7 f orbitals 16 32 Chapter 5.3 Vocabulary • • • • • • 61 Electron configuration Aufbau principle Pauli Exclusion Principle Hund’s Rule Valence Electron Electron Dot Structure •Electron Configurations • An atom’s electron configuration is the arrangements of electrons in the atom. • Electrons are arranged to minimize energy. • In other words, electrons fill up the lowest energies possible first. This is the Aufbau Principle. 62 •Assigning Electrons to Atoms • Electrons generally assigned to orbitals of successively higher energy. • For • For H atoms, E depends only on n. many-electron atoms, energy depends on both n and l. 63 •Energy Level Diagram of Hydrogen 64 •Assigning Electrons to Subshells In H atom all subshells of same n have same energy. In many-electron atom: a) subshells increase in energy as value of n + l increases. b) for subshells of same n + l, subshell with lower n is lower in energy. • 65 Orbitals and Their Energies Many-Electron Atoms 66 •Aufbau Diagram -- Filling Electron Orbitals 1s 2s 3s 4s 5s 6s 7s 3p 4p 5p 6p 3d 4d 5d 8s Start here n+l=1 n+l=2 n+l=3 2p n+l=4 n+l=5 n+l=6 4f n+l=7 6d 7p Haven’t gotten this far. What orbitals are being filled with elements 110-118? 5f n+l=8 The orbital with the lower ‘n’ is lower in energy if the n+l number is the same. 67 •Writing Atomic Electron Configurations Two ways of writing configs. One is called the spdf notation. spdf notation for H, atomic number = 1 1 1s value of n 68 no. of electrons value of l •Pauli Exclusion Principle No two electrons in the same orbital can have the same spin. One electron is spin up, the other is spin down. 69 •Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. ORBITAL BOX NOTATION for He, atomic number = 2 Arrows 2 depict electron spin 1s 1s It would be a violation of the Pauli exclusion principle to have both of these electrons as spin up or both as spin down. 70 •Electron Configurations and the Periodic Table 71 •Lithium Group 1 (1A) Atomic number = 3 1s22s1 3 total electrons 3p 3s 2p 2s 1s 72 Interactive Periodic Table Ground State Electron Configurations 73 •Carbon Group 14 (4A) Atomic number = 6 1s2 2s2 2p2 6 total electrons 3p 3s 2p 2s 1s 74 Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. Electron Configuration for Elements 11-18 Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements. 75 Sodium Group 1 (1A) Atomic number = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 [Ne] 3s1 (uses noble gas notation) Note that we have begun a new period. All Group 1A elements have [core]ns1 configurations. 76 Electron Configurations and the Periodic Table 77 Transition Metals All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are d-block elements. Chromium 78 Iron Copper 79 Electrons in Energy Levels • Electrons fill up levels from lowest energy to highest energy (Aufbau Principle) • Outermost electrons are called Valence Electrons. • When atoms come close together it is the Valence Electrons that interact. 80 Valence Electrons • How to determine which electrons are in outer shell? • Write down electron configuration of an element in noble-gas configuration • Whatever electrons are displayed in the highest energy shell (n) only are valence electrons (Main Group elements) 81 Lewis Dot Diagrams • How do we represent Valence Electrons? • By a Lewis Dot Diagram • Rules for Lewis Dot Diagrams: • Use the Elemental Symbol • Use 1 dot to represent each valence electron • The symbol represents the nucleus and all the inner (core) electrons. • Examples: . .. .. • Li·, Be: , ·C·, . ·Cl:, .. :Ne: .. 82 Valence Electrons Examples • O given by [He]2s22p4 so it has 6 valence electrons. • • 83 Ga given by [Ar]3d104s24p1 has 3 valence electrons. •Electron Dot Representation 84